1 / 26

Acids, Bases & Salts

Acids, Bases & Salts. Chapter 19 Pages 586-629. 19.1 Acid-Base Theories. Acids - taste sour, will change the color of an acid-base indicator (red), and can be strong or weak electrolytes in aqueous solutions Common acids: vinegar, carbonated drinks, citrus foods. Common Acids & their Names.

olympia-clarke
Télécharger la présentation

Acids, Bases & Salts

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Acids, Bases & Salts Chapter 19 Pages 586-629

  2. 19.1 Acid-Base Theories • Acids - taste sour, will change the color of an acid-base indicator (red), and can be strong or weak electrolytes in aqueous solutions • Common acids: vinegar, carbonated drinks, citrus foods

  3. Common Acids & their Names • HCl - hydrochloric acid • HNO3 - Nitric acid • H2SO4 - sulfuric acid • H3PO4 - Phosphoric acid • CH3COOH - ethanoic acid (vinegar) • H2CO3 - Carbonic acid

  4. Bases • Defined by their bitter taste, slippery feeling; will change the color of an acid-base indicator (blue color), can be strong or weak electrolytes in aqueous solution • Soap, salt water, milk of magnesia

  5. Arrhenius Acids & Bases • Arrhenius gave a more specific definition to acids and bases aside from their feel and taste. • He said that acids are hydrogen-containing compounds that ionize to yield H+ in aqueous solutions. • He said that bases are compounds that ionize to yield OH- in aqueous solutions

  6. Classification of acids • HCl - hydrochloric acid (monoprotic) • HNO3 - Nitric acid (monoprotic) • H2SO4 - sulfuric acid (diprotic) • H3PO4 - Phosphoric acid (triprotic) • CH3COOH - ethanoic acid (vinegar) (triprotic) • H2CO3 - Carbonic acid (diprotic) • You write the definition of mono-, di- and triprotic

  7. Bronsted-Lowry Acids & Bases • Used because Arrhenius isn’t very detailed • Defines acids as H+ donors and a base as a H+ acceptor

  8. Conjugate Acids & Bases • NH3 + H2O --> NH4+ + OH- • base + acid --> conj acid + conj base

  9. Lewis Acids & Bases • Acids accept a pair of electrons during a reaction • Bases donate a pair of electrons • H+ + OH- --> H2O • Lewis acid + Lewis base --> cmpd

  10. 19.2 Hydrogen Ions & Acidy • Inside a container of water, occasional collisions between water compounds occur with enough energy to form ions. • HOH + HOH  H3O+ + OH- • H3O+ are known as hydronium ions (can be simplified to H+) • This happens in small occasions and will occur in equal concentrations of ions.

  11. Concentration of Ions • Water acts as the neutral solution when calculating the pH of any solution. • The concentration (mol/L) of the ions (H3O+ and OH-) are each said to be 1 x 10-7 M. • The overall concentration of water is then calculated [H3O+] x [OH-] = 1 x 10-14 M= Kw • This is called the ion-product constant for water – Kw.

  12. Acidic Concentrations • Acidic solutions are when the [H+] is greater than the [OH-]. • The [H+] will be greater than 1 x 10-7 M. • The [OH-] will be less than 1 x 10-7 M. • Ex: HCl  H+ + Cl-

  13. Basic Concentrations • Basic solutions are when the [H+] is less than the [OH-]. (Are also called alkaline solutions.) • The [H+] will be less than 1 x 10-7 M. • The [OH-] will be greater than 1 x 10-7 M. • Ex: NaOH  Na+ + OH-

  14. Calculating the pH value • pH values range from 0 – 14 where 0 is extremely acidic, 7 is neutral and 14 is extremely basic • pH = -log[H+] • (Log is a button on your calculator) • pOH = -log[OH-] • pOH + pH = 14

  15. Calculations • Using [H+] to solve for pH • Using [OH-] to solve for pOH • Using [OH-] to solve for pH • Using [H+] to solve for pOH • Using pH to solve for [H+] • Using pOH to solve for [OH-]

  16. Acid-Base Indicators • Solutions can cause litmus papers to change color indicating acid/base. • Indicators can be added to a solution to produce a color change of the solution at specific pH values. • pH meters can register an accurate measurement of a pH value.

  17. 19.3 Strengths of Acids & Bases • Strength of acids or bases depend on how well the compound ionizes in water. • Strong Acids - completely ionize in water (totally break apart creating a charged particle in the water) HCl • Weak Acids - ionize only slightly (may still have a H-ion that could be removed)

  18. Equilibrium-Constant Expression

  19. Acid Dissociation Constant • Tells whether an acid is strong or weak - stronger acids have higher values.

  20. Sample problem • A 0.1000 M solution of ethanoic acid is only partially ionized. From measurements of the pH of the solution, [H+] is determined to be 1.34 x 10-3 M. What is the Ka?

  21. Solving the problem • [H3O+] = 1.34 X 10-3 M • [ACID] = 0.1000 M • [A-] = (0.1000 - 1.34 X 10-3 M) = 0.0987

  22. Base Dissociation Constant • Strong bases - dissociate completely into metal ions and OH- ions. • Weak bases - dissociate into the conjugate bases and OH- ions.

  23. 19.4 Neutralization Reactions • Acid-base Reactions • Mixing strong acids w/ strong bases will produce a salt and water. • The final product will be neutral as long as the correct mole:mole relationship is followed. • HCl + NaOH --> NaCl + H2O

  24. Titration • Using a buret to add acid to a base to find the equivalence point (where a color change occurs) • The will allow the [base] to be determined because the [A] is already known. • Titration

  25. Steps of Titration • A known [A] is added to a flask (standard solution) • Several drops of an indicator are added • Measured values of base are mixed with the acid until a slight color change is noticed. (end point)

  26. Calculating unknown [B] A 25-mL soln of H2SO4 is completely neutralized by 18 mL of 1.0 M NaOH. What is the concentration of the H2SO4 soln?

More Related